Micellar Effects on the Electron Transfer Reaction within the Ion Pair

by potassium ferricyanide in the presence of N,N-dimethyldodecylamine N-oxide. Shehata M. Zourab , Essam M. Ezzo , Hisham J. El-Aila , Jamil K. J...
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16978

J. Phys. Chem. 1996, 100, 16978-16983

Micellar Effects on the Electron Transfer Reaction within the Ion Pair [(NH3)5Co(N-cyanopiperidine)]3+/[Fe(CN)6]4Amalia Rodrı´guez, Marı´a del Mar Graciani, Robert Balahura,† and Marı´a Luisa Moya´ * Departamento de Quı´mica Fı´sica, UniVersidad de SeVilla, c/o Professor Garcı´a Gonza´ lez s/n, 41012 SeVilla, Spain, and Department of Chemistry and Biochemistry, UniVersity of Guelph, Guelph, Ontario N162W1, Canada ReceiVed: April 24, 1996; In Final Form: August 1, 1996X

The electron transfer reaction between [Co(NH3)5(N-cyanopiperidine)]3+, pentaammine(N-cyanopiperidine)cobalt(III), and hexacyanoferrate(II) has been studied in aqueous solutions. Typical saturation behavior was observed at high hexacyanoferrate(II) concentrations, which allows the precursor complex formation equilibrium constant, KIP, and the true electron transfer rate constant, ket, to be obtained. Variations of ket in the presence of various background electrolytes (sodium sulfate and sodium chloride) and surfactants (sodium dodecyl sulfate, SDS, dodecyltricosaethylene glycol ether, Brij35, and octylphenol(ethylene oxide)9.5 ether, Triton X-100) have been investigated at 298.2 K. Rationalization of the experimental data was assisted by the use of the Marcus theory on electron transfer. The conclusion is that micellar effects on the electron transfer rate constant can be explained by considering the micelles as a special background electrolyte with a high charge and a strong power of hydration.

Introduction Traditionally, variations of the rate constant of an overall electron transfer process with surfactant concentration have been treated quantitatively in terms of the pseudophase model, which includes distribution of one or both of the reactants between water and micelles.1 As an example, for a bimolecular reaction, in the case that the reactant present in large excess, A, is an anion in an anionic micellar system (and the other reactant, B, is adsorbed at the micellar surface), the observed pseudo-firstorder rate constant for the overall reaction is given by1

k)

kw[Aw] + k2mKs[Dn][A]∆ 1 + Ks[Dn]

(1)

Here kw and k2m are second-order rate constants for the reaction in the aqueous and micellar pseudophases respectively, Ks is the equilibrium binding constant of the species B to the micelles, Dn is the concentration of the micellized surfactant, and [Aw] and [A]∆ are the molarity of species A in water and in a shell of thickness ∆ at the micellar surface, respectively. [A]∆ can be estimated by solving the Poisson-Boltzmann equation in spherical symmetry.2 To explain micellar effects on k according to eq 1, it is necessary to calculate k2m and therefore [A]∆. Unfortunately, the calculated values of [A]∆ depend (among other factors, the size of the micelles, the aggregation number, ...) on the ∆ value that is used to perform the calculations, and so does k2m.3 This approach explains the experimental changes observed in k through the changes in [A]∆ resulting from changing surfactant concentration. This is a simplification of the problem since other factors can also change with surfactant concentration, thus contributing to the changes in ket (the use of eq 1 involves the implicit assumption that Ks, kw, and k2m are unaffected by changes in [surfactant] or other species concentrations in the aqueous medium). * Author to whom all correspondence should be directed. † University of Guelph. X Abstract published in AdVance ACS Abstracts, October 1, 1996.

S0022-3654(96)01175-6 CCC: $12.00

A way to circumvent the above-mentioned problem is to use the classical formulation for electron transfer reactions and then consider micellar effects upon the different parameters that determine the rate constant. An outer-sphere electron transfer reaction can be viewed as follows:

A + B a A/D A/D f A-/D+ V products

KIP ) k1/k-1 ket

where KIP is the precursor complex formation equilibrium constant and ket is the rate constant of the electron transfer step. Depending on the magnitudes of k1, k-1, and ket the observed rate constant can be expressed with various equations.4 As an example of two limiting cases, if ket . k-1, that is, when step 1 is a rapid preequilibrium and the second step is the slow electron transfer from the donor to the acceptor, kobs ) KIPket. If ket , k-1 then kobs ) k1. The micellar system can influence ket through each of the terms which contribute to its magnitude. On the basis of the Marcus theory5 these terms are the reorganization energies, the driving force accompanying the reaction, the electronic transmission coefficient, and the nuclear frequency factor (see below). In regard to KIP, changes in surfactant concentration can affect KIP when the reactants are ions as well as in the case when they are not. In the first case it is expected that Coulombic contributions are the more important factors to consider, whereas in the second case other type of contributions are operative. From the above comments it is possible to see that kinetic micellar effects on an electron transfer process can be the sum of many contributions to the experimentally observed rate constant. On this basis, and in order to simplify the study of micellar effects, it is appropriate to study a system for which it is possible to follow the true electron transfer rate constant, ket, independently of KIP. With this idea in mind, the reaction between hexacyanoferrate(II) and pentaammine(N-cyanopip© 1996 American Chemical Society

The Ion Pair [(NH3)5Co(N-cyanopiperidine)3+/[Fe(CN)64-] eridine)cobalt(III) (this cobalt(III) complex is shown in structure 1) was chosen. From the high and opposite charges of the 3+

J. Phys. Chem., Vol. 100, No. 42, 1996 16979

SCHEME 1 KIP

[Co(NH3)5(N-cyanopiperidine)]3+ + [Fe(CN)6]4- {\} [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4-

(NH3)5Co NC N 1

reactants one can expect a large KIP value, that is, the possibility of observing saturation behavior (see below) and, therefore, of obtaining the ket value. In the present work, this reaction has been studied in aqueous solutions in the presence of sodium sulfate and sodium chloride as well as in micellar solutions of sodium dodecyl sulfate, SDS, dodecyl tricosaoxyethylene glycol ether, Brij35, and octylphenol(ethylene oxide)9.5 ether, Triton X-100. In all the experiments temperature was kept at 298.2 K. Experimental Section Materials. Pentaammine(N-cyanopiperidine)cobalt(III) perchlorate was prepared by a method previously described.6 Caution! Perchlorate salts are potentially explosive and should be handled with care. The complex was purified by cation exchange chromatography on CM-Sephadex. Sodium hexacyanoferrate(II), sodium chloride, sodium sulfate, and disodium edta were Merck P.A. grade. Sodium dodecyl sulfate and Brij35 were also from Merck, and Triton X-100 was kindly supplied by Professor M. A. De la Rosa. Surfactants were used as purchased. Double-distilled water was used. Kinetics. The rate measurements were performed using a Hitachi 150-20 UV-visible spectrophotometer. When the reaction time was less than 15 min, a manual mixing apparatus from Hi-Tech was used. The kinetics of the reaction studied, [Co(NH3)5(N-cyanopiperidine)]3+ + [Fe(CN)6]4-, was followed at 420 nm, the wavelength of the maximum of the hexacyanoferrate(III) ion produced in the process. All solutions were freshly prepared and N2 was bubbled through the water before preparing the micellar solutions to avoid oxidation of the hexacyanoferrate(II) ions. The cobalt(III) complex concentration in all the reaction media was 1.5 × 10-4 mol dm-3. All runs employed an excess of hexacyanoferrate(II). Na2(H2edta) was added to prevent precipitation of Co2+ ions produced in the reaction. According to preliminary experiments the reaction rate was not affected by changing the Na2(H2edta) concentration from 1.5 × 10-4 to 5 × 10-4 mol dm-3. In this work, the higher concentration was used. The observed rate constant was obtained from the slopes of the ln(A∞ - At) against time plots, where A∞ and At are the absorbances at the end of the reaction and at a time t. The plots were good straight lines for at least three half-lives. The precision in the rate constants was within 5%. pH Measurements. The pH measurements were performed on a micropH 2000 Crison instrument. In all cases the pH of the working solutions was >6 and therefore the unique hexacyanoferrate(II) species present in the reaction medium is the ion [Fe(CN)6]4- ([Fe(CN)6H]3- a [Fe(CN)6]4- + H+, pK ) 4.327).

ket

[Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- 98 [Co(NH3)5(N-cyanopiperidine)]2+/[Fe(CN)6]3- f products [Fe(CN)64-].[Co(III)] the observed rate constant can be written as

kobs )

ketKIP[Fe(CN)64-] 1 + KIP[Fe(CN)64-]

(2)

where KIP is the equilibrium constant for the precursor complex formation step and ket is the electron transfer rate constant. On the basis of eq 2 and in the presence of high hexacyanoferrate(II) concentration, one would expect saturation behavior, that is, the possibility of reaching a plateau in which the observed rate constant is independent of the hexacyanoferrate(II) concentration and therefore equal to ket. Figure 1 shows the plot of kobs against [Fe(CN)64-] at constant ionic strength (I ) 0.5 mol dm-3, the ionic strength was kept constant by adding the requisite amount of NaCl in each case). Saturation is reached at [Fe(CN)64-] ) 0.02 mol dm-3 and ket ) (3.08 ( 0.04) × 10-2 s-1. The KIP value can be obtained from the plot of kobs-1 against [Fe(CN)64-]-1. This plot is a good straight line, as expected from eq 2, and gives KIP ) 420 ( 20 mol-1 dm3. This value is larger than that expected only on the basis of Coulombic interactions at 0.5 mol dm-3 ionic strength (∼50 mol-1 dm3), showing the importance of the hydrogen-bonding interactions within the ion pair, as has been suggested for similar [amminecobalt(III)]3+/[Fe(CN)6]4- ion pairs.8 Table 1 shows the ket values in sodium chloride and sodium sulfate aqueous solutions (these ket values have been determined from the values of kobs at the saturation level). Saturation was checked for the higher electrolyte concentration used in each case (by investigating the influence of hexacyanoferrate(II) concentration on the observed first-order rate constant), to be sure that the observed rate constant was equal to ket. In all cases for [Fe(CN)64-] ) 0.03 mol dm-3, saturation was reached. Figure 2a shows the plot of kobs against [Fe(CN)64-] for [SDS] ) 0.25 mol dm-3. At 0.03 mol dm-3 of hexacyanoferrate(II) ions concentration, saturation behavior was reached. A wider range of hexacyanoferrate(II) concentration could not be studied

Results The reaction of [Co(NH3)5(N-cyanopiperidine)]3+ with [Fe(CN)6]4- can be written as in Scheme 1. Under the condition

Figure 1. Plot of kobs against hexacyanoferrate(II) concentration for the reaction [Fe(CN)6]4- + [Co(NH3)5(N-cyanopiperidine)]3+ in aqueous solution at 298.2 K (I ) 0.5 mol dm-3 NaCl).

16980 J. Phys. Chem., Vol. 100, No. 42, 1996

Rodrı´guez et al.

TABLE 1: Dependence of the Observed Rate Constant on Hexacyanoferrate(II) Concentration (I ) 0.5 mol dm-3 NaCl) for the Reaction [Co(NH3)5(N-cyanopiperidine)]3+ + [Fe(CN)6]4-, kobs/s-1, and First-Order Rate Constants, ket (s-1), for the Electron Transfer Reaction within the Ion Pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- in Sodium Chloride and Sodium Sulfate Aqueous Solutions at 298.2 K 103[Fe(CN)64-]/ mol dm-3

102kobs/s-1

103[Fe(CN)64-]/ mol dm-3

102kobs/s-1

2.53 2.53 5.00 6.34 7.53

1.70 1.70 2.35 2.67 2.75

10.0 13.1 16.3 20.0 25.0

2.91 3.01 3.08 3.09 3.08

0.100 0.150 3.5 3.2 0.040 0.065 0.075 3.0 2.5 2.3

0.200 3.0 0.100 0.109 2.1 2.0

[NaCl]/mol dm-3 102ket/s-1 a [Na2SO4]/mol dm-3 102ket/s-1 a a

0.000 3.8 0.000 3.8

0.050 3.6 0.0218 3.4

[Fe(CN)64-] ) 0.03 mol dm-3.

Figure 3. Plot of kobs against hexacyanoferrate(II) concentration for the reaction [Fe(CN)6]4- + [Co(NH3)5(N-cyanopiperidine)]3+ in (a) [Triton X - 100] ) 0.01 mol dm-3 and (b) [Brij35] ) 0.01 mol dm-3 solutions at 298.2 K (I ) 0.5 mol dm-3 NaCl).

Figure 2. Plot of kobs against hexacyanoferrate(II) concentration for the reaction [Fe(CN)6]4- + [Co(NH3)5(N-cyanopiperidine)]3+ in (a) [SDS] ) 0.25 mol dm-3 and (b) [SDS] ) 0.25 mol dm-3 + 0.2 mol dm-3 NaCl solutions at 298.2 K (ionic strength was kept constant by adding the requisite amount of NaCl in each case).

because of the ionic strength limitation (0.5 mol dm-3) and therefore the observed rate constant was not comparable to the others. In order to work within a wider hexacyanoferrate(II) concentration range, the reaction was also studied in SDS solutions but in the presence of 0.2 mol dm-3 of NaCl. Figure 2b shows the plot of kobs against [Fe(CN)64-] for [SDS] ) 0.25 mol dm-3 in the presence of NaCl 0.2 mol dm-3. This figure shows that at [Fe(CN)64-] ) 0.03 mol dm-3 saturation was reached. From the results obtained from the study of the dependence of kobs on hexacyanoferrate(II) concentration in SDS solutions, in the absence and in the presence of 0.2 mol dm-3 of NaCl, one can say that the experimentally observed rate constant is equal to ket when [Fe(CN)64-] ) 0.03 mol dm-3. Therefore, the kinetic data summarized in Table 2 correspond to the true electron transfer rate constant, ket, which have been determined from the values of kobs at the saturation level. In

regard to the data shown in Table 2, kinetic experiments could not be carried out at SDS concentrations lower than 0.04 mol dm-3 due to solubility problems in relation to the reactant Co(III) complex. Plots of kobs against [Fe(CN)64-] for two Brij35 and Triton X-100 solutions appear in Figure 3. In both cases the ionic strength was kept constant with sodium chloride. One can see from these figures that at [Fe(CN)64-] ) 0.03 mol dm-3 saturation level has been reached. Table 3 shows the ket values for the reaction studied in nonionic surfactant solutions. They have been determined, as before, from the observed rate constant values at the saturation level. Equation 2 fits well the experimental variations of kobs with hexacyanoferrate(II) concentration shown in Figures 1-3. In relation to the critical micelle concentration (cmc) of the different micellar systems used, for the SDS micelles a cmc < 8 × 10-3 mol dm-3 is expected in all cases, given the high ionic concentration present in the reaction medium.9a For this anionic surfactant the measurement of the cmc in the presence of the Co(III) reactant was not possible because of solubility problems at low surfactant concentration. In regard to the Brij35 and Triton X-100, the former shows a cmc of 6 × 10-5 mol dm-3 9b and the latter a cmc of 2.7 × 10-4 mol dm-3,9c both in aqueous solutions. As will be seen below, the cmc values are not used in the discussion of the kinetic data. Discussion From the examination of the data listed in Tables 1 and 2, one can see the strong kinetic effect of the SDS micelles on the electron transfer within the ion pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- (kobs ) ket). Before analysis of the energetic terms which can be responsible for the observed micellar effects, it is important to assess what changes are expected to be produced due to the micellar systems used as

The Ion Pair [(NH3)5Co(N-cyanopiperidine)3+/[Fe(CN)64-]

J. Phys. Chem., Vol. 100, No. 42, 1996 16981

TABLE 2: First-Order Rate Constants, ket (s-1), for the Electron Transfer Reaction within the Ion Pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- in Sodium Dodecyl Sulfate Solutions at 298.2 Ka [SDS]/mol dm-3 b 104ket/s-1 [SDS]/mol dm-3 c 104ket/s-1 a

0.040 11.2 0.040 29

0.050 9.5 0.050 25

0.075 8.2 0.075 18

0.100 7.3

0.125 6.4 0.100 16

0.15 5.8 0.150 12.2

0.175 5.6 0.200 10.3

0.200 5.8 0.225 9.4

0.225 6.6

0.250 8.4 0.250 8.4

[Fe(CN)64-] ) 0.03 mol dm-3. b SDS. c SDS + 0.2 mol dm-3 NaCl.

TABLE 3: First-Order Rate Constants, ket(s-1), for the Electron Transfer Reaction within the Ion Pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- in Nonionic Surfactants at 298.2 K (0.2 mol dm-3 NaCl)a [Brij35]/mol dm-3 102ket/s-1 [Triton X-100]/mol dm-3 102ket/s-1 a

10-4 4.1

10-3 4.0 2.5 × 10-3 3.1

5 × 10-3 4.3

10-2 4.1 5 × 10-3 3.3

2.5 × 10-2 4.2 10-2 3.1

5 × 10-2 4.2 t × 10-2 3.2

7.5 × 10-2 4.3 7.5 × 10-2 3.1

[Fe(CN)64-] ) 0.03 mol dm-3.

reaction media by changing surfactant concentration. It has been pointed out that in the presence of high ionic concentrations, variations in the [SDS] do not cause substantial changes in the micellar characteristics (aggregation number, dissociation degree, ...10). This has also been shown by the experimental data obtained by Pramauro et al.11 for the reaction Fe2+ + IrCl62in SDS micelles. Nonetheless, in spite of the characteristics of the micelles not changing, the ionic concentration in the reaction medium increases by increasing SDS concentration. With this in mind and with the idea of investigating the influence of the ionic concentration on the electron transfer within the ion pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- in a more simple reaction medium, this process has been studied in aqueous electrolyte solution (Table 1). Experimental results show that ket decreases as the ionic strength of the medium increases, sodium sulfate being more effective in retarding the reaction than sodium chloride at a given ionic strength. To rationalize data in aqueous salt solutions, the Marcus theory was used to express the electron transfer rate constant as5

ket ) κelνn exp(-∆Gq/RT)

(3)

Here κel is the electronic transmission coefficient and νn is the nuclear frequency factor. ∆Gq is the free energy of activation. From data in the literature concerning similar electron transfer reactions,12 the process under study in this work can be considered adiabatic and therefore κel ) 1. On the other hand, for an adiabatic reaction, for which the Ochinikova condition is not fulfilled13 νn depends only on the internal vibrational modes14 and considering that there is no coupling between internal and external (solvent) modes, νn can be assumed to be constant when changing the reaction medium properties. In regard to the free energy of activation, it can be written5

∆Gq )

λ ∆G°′ 2 1+ 4 λ

(

)

(4)

where the reorganization energy, λ, is the sume of λin and λout, the first term being the contribution of the intramolecular reorganization energy related to changes in the metal-ligand bonds during the activation process (assumed to be constant with changes in the reaction medium) and the second term the sum λout ) λsol + λatm, where λsol is the contribution of the solvent reorganization energy and λatm is the ionic cloud reorganization energy. ∆G°′ is the free energy change accompanying the electron transfer within the ion pair. Taking into account the above comments, the influence of the ionic strength on ket is expected to operate through λatm,

λsol, and ∆G°′. The solvent reorganization energy, in its classical expression, can be written as5

λsol ) Ne2

(

)(

1 1 1 + a1 a2 R

1 1 Dop Ds

)

(5)

Here a1, a2, and R are the reactant metal center radii and the closest approximation distance between the metal centers during the reaction, respectively. Dop and Ds are the optical and static dielectric constants of the reaction medium. In salt solutions of the concentrations used in this work, the Pekar’s factor (1/ Dop - 1/Ds), changes slightly and so does λsol. In respect to the ionic cloud reorganization energy contribution, it can be calculated by means of the mean spherical approximation (MSA).15 This contribution from a change in ionic strength from 0 to 1 mol dm-3 is smaller than +0.2 kJ mol-1.16 From these considerations λout is not expected to change substantially, especially if one takes into account that not only λsol and λatm show small changes by changing ionic strength in the concentration range used in this work, but the signs of these contributions are opposite. Therefore, the expected changes in λout will be even smaller than those shown by λsol and λatm, respectively. All this means that in the present case, the main factor responsible for the electron transfer rate constant changes by changing ionic strength will be ∆G°′. In regard to the changes in ∆G°′ with salt concentration for the Fe(CN)63-/4- couple, experimental data show that the oxidation of hexacyanoferrate(II) is more difficult when the concentration of background electrolytes in the medium increases.17 Besides, there is a specificity of these salt effects: the higher the charges of the ions from the background electrolytes, the greater the stabilization of the hexacyanoferrate(II) ions. For the cobalt(III)/(II) couple there are no experimental data since this couple is irreversible from the electrochemical point of view due to the lability of the Co(II) complex. Nonetheless, on the basis of a primitive model, it would be expected that the reduction of the Co(III) center will be more difficult when the salt concentration in the medium increases as the charge of the Co(III) complex is higher than that of Co(II). Therefore, the term ∆G°′ would be expected to retard the reaction with increasing salt concentration; that is, a decrease in ket with increasing [salt] is expected and observed. Now kinetic data in micellar solutions will be considered. It is worth noting that the reactant species is the ion pair (under the working conditions kobs ) ket) which has, as a whole, a negative charge. Therefore, this species is not expected to be located within the negative Stern layer of the micelles and the process will take place in a zone of the system in which the ionic strength is low enough to permit saturation behavior to be reached. Taking this into account, and for the same reasons

16982 J. Phys. Chem., Vol. 100, No. 42, 1996

Figure 4. Plot of ket against sodium chloride concentration for the electron transfer reaction within the ion pair [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- in [SDS] ) 0.1 mol dm-3 at 298.2 K.

mentioned above in relation to the salt effects in conventional aqueous solutions, small changes in λout are expected by varying [SDS] and therefore the changes in ∆G°′ will also control the dependence of ket on the SDS concentration. The changes in ∆G°′ can be calculated through activity coefficient variations. These activity coefficients can be expressed as a function of the ionic strength present in the reaction medium, but their magnitude is not known in our case. Nonetheless, in a qualitative way, the trends observed in ket with SDS concentration can be explained. In the presence of SDS + 0.2 mol dm-3 NaCl, ket decreases when the surfactant concentration increases. For SDS systems at low surfactant concentration the same trend was found, but the electron transfer rate constant reaches a minimum and later on ket increases by increasing SDS concentration. Since the different behavior in the two cases is due to the presence of sodium chloride, it seems interesting to see what changes in the micellar systems are caused by the presence of this salt. Published results show that when the amount of sodium chloride increases, the aggregation number, Nag, of the SDS micelles increases.18 As an example, for 0.3 mol dm-3 NaCl Nag ∼ 110, whereas for NaCl 0.5 mol dm-3 Nag ∼ 300. That is, the micelle concentration decreases by increasing [NaCl]. On the other hand, the electrical potential at the SDS micelle surface becomes less negative when NaCl concentration increases.19,20 If the micelles are considered as a kind of electrolyte, an increase in [NaCl] would lead to a decrease in the concentration of the micellar electrolyte as well as a decrease in its effectiVe charge. With this in mind, and considering that variations in ∆G°′ are the main factor controlling reactivity, an increase in ket would be expected by increasing NaCl concentration at constant [SDS]. Figure 4 shows the variation of ket with [NaCl] at [SDS] ) 0.1 mol dm-3. This figure confirms the expected trend. Traditionally, kinetic salt effects in reactions between ions have been explained by taking into account: (i) Coulombic interactions between the participants in the reaction and the ions from the added salt and (ii) non-Coulombic interactions between the ions from the background electrolyte and the solvent. These non-Coulombic interactions, in terms of a hydration model,21 result in the diminution of the free water molecules, since water bound to ions (water molecules in the hydration shells of the ions from the added salt) is no longer a part of the bulk solvent. Coulombic interactions stabilize ions when the electrolyte concentration increases. Non-Coulombic interactions are responsible for the destabilization of ions as electrolyte concentration increases. At low ionic strength the Coulombic interactions are the main factors operating on reactivity. As ionic strength

Rodrı´guez et al. increases (because electrolyte concentration increases or because there is an increase in its charge) the importance of the nonCoulombic interactions also increases, becoming the most important contributions at high ionic strength.22 Of course, the range of electrolyte concentration for which one or the other is dominant depends on the nature of the background electrolyte (charge, size, ...). In the present case, for the electron transfer reaction within the ion pair [Co(NH3)5(N-cyanopiperidine)]3+/ [Fe(CN)6]4-, Coulombic interactions would cause a greater stabilization of the initial state in respect to the less polar transition state, retarding the reaction as the concentration of added salt increases. On the contrary, non-Coulombic interactions would provoke a greater destabilization of the initial state than of the transition state as the electrolyte concentration increases, accelerating the reaction. Given that the SDS micelles in the absence of 0.2 mol dm-3 NaCl can be considered a stronger electrolyte than in the presence of this salt, it would be expected that the non-Coulombic effect could become dominant at lower SDS concentrations since micellar concentration, that is, electrolyte concentration, is higher and its effective charge seems to be higher too. One would thus expect the appearance of a minimum in the case of SDS micellar systems at lower SDS concentrations compared to SDS + 0.2 mol dm-3 NaCl micellar systems, as in fact is observed. Kinetic data seem to indicate that ket values also tend to reach a minimum in the latter case. Unfortunately, higher concentrations of SDS could not be used because of solubility problems. In regard to the results obtained in Brij35 and Triton X-100 micellar solutions in the presence of 0.2 mol dm-3 NaCl, in both cases a small increase in ket is found with respect to aqueous solutions. No dependence of ket on surfactant concentration was found. Considering the ionic character of the reactant ion pair, no strong interactions are expected with neutral micelles, and, of course, there will not be changes in ionic concentration by changing [surfactant]. The slight increase found in ket with respect to aqueous solutions could be related to a hydrophobic effect. When the reaction is studied in aqueous solutions (under the same working conditions) in the presence of 0.2 mol dm-3 of (Et)4NNO3 (tetraethylammonium nitrate) ket ) 5.3 × 10-2 s-1 (compared to 3.05 × 10-2 s-1 in the presence of 0.2 mol dm-3 NaCl). A further increase in the Et4NNO3 concentration causes an increase in ket. This observed trend is the opposite to that found for the dependence of ket on salt concentration for typical hydrophilic salts. The same opposite trends were found for NaCl, NaNO3, etc., which originate a hypsochromic shift, and hydrophobic electrolytes, which originate a bathochromic shift, in the study of the metal-to-ligand charge transfer bands of some iron(II) complexes.23 This result has been related to a stabilization of the less polar species in the presence of these hydrophobic salts. To summarize, micellar effects in SDS solutions on the [Co(NH3)5(N-cyanopiperidine)]3+/[Fe(CN)6]4- f [Co(NH3)5(Ncyanopiperidine)]2+/[Fe(CN)6]3- reaction can be rationalized by noting that changes in ∆G°′ are the main factor affecting reactivity and considering that SDS micelles behave as special background electrolytes, with a higher charge and a stronger power of hydration than the usual salts. Acknowledgment. The authors acknowledge financial support of the D.G.I.C.Y.T (PB-92-0677) and Consejerı´a de Educacio´n y Ciencia de la Junta de Andalucı´a. References and Notes (1) Gra¨tzel, M. Micellization, Solubilization, and Microemulsions; Mital, K. L., Ed.; Plenum Press: New York, 1977; Vol. 2.

The Ion Pair [(NH3)5Co(N-cyanopiperidine)3+/[Fe(CN)64-] (2) (a) Rodenas, E.; Ortega, F. J. Phys. Chem. 1987, 91, 837. (b) Bunton, C. A.; Moffatt, J. J. Phys. Chem. 1988, 92, 2896. (3) Table VI of ref 2a. (4) Sa´nchez-Burgos, F.; Moya´, M. L.; Gala´n, M. Prog. React. Kinet. 1994, 19, 1. (5) Cannon, R. D. In Electron Transfer Reactions; Butterworths: London, 1980. (6) Alexander, C. S.; Balahura, R. J. Inorg. Chem. 1994, 33, 1399. (7) Kershaw, M. R.; Prue, J. R. Trans. Faraday Soc. 1967, 63, 1198. (8) (a) Miralles, A. J.; Armstrong, R. E.; Haim, A. J. Am. Chem. Soc. 1977, 99, 1416. (b) Szecsy, A. P.; Haim, A. J. Am. Chem. Soc. 1981, 103, 1679. (9) Van Os, N. M.; Haak, J. R.; Rupert, L. A. M. Physico-Chemical Properties of Selected Anionic, Cationic and Nonionic Surfactants; Elsevier: Amsterdam, 1993; (a) p 24; (b) p 218; (c) p 310. (10) Frahm, J.; Dieckman, S.; Haase, A. Ber. Bunsenges Phys. Chem. 1980, 80, 566. (11) Pramauro, E.; Pelizzetti, E.; Dieckman, S.; Frahm, J. Inorg. Chem. 1982, 21, 2432. (12) Pe´rez-Tejeda, P.; Benko, J.; Moya´, M. L.; Sa´nchez, F. J. Mol. Liq. 1995, 65/66, 261. (13) (a) Doine, H.; Swaddle, T. W. Can. J. Chem. 1988, 66, 2769. (b) Grampp, G.; Harrer, W.; Jaenike, W. J. Chem. Soc., Daraday Trans. 1 1987, 83, 161.

J. Phys. Chem., Vol. 100, No. 42, 1996 16983 (14) Sumi, H.; Marcus, R. A. J. Chem. Phys. 1986, 84, 4894. (15) German, E. D.; Kuznetsov, A. M. Electrochimiya 1987, 23, 1560. (16) Gala´n, M.; Domı´nguez, M.; Andreu, R.; Moya´, M. L.; Sa´nchez, F.; Burgess, J. J. Chem. Soc., Faraday Trans. 1990, 86, 937. (17) Gala´n, M.; Jime´nez, R.; Sa´nchez, F. Ber. Bunsenges Phys. Chem. 1993, 93, 16. (18) Mazer, N. A.; Benedek, G. B.; Carey, M. C. J. Phys. Chem. 1976, 80, 1075. (19) Fe´rnandez, M. S.; Fromherz, P. J. Phys. Chem. 1977, 81, 1755. (20) Sjo¨blom, J.; Gestblom, B. In Organized Solutions; Friberg, S. E., Lindman, B., Eds.; Marcel Dekker: New York, 1992; p 193. (21) (a) Robinson, R. A.; Stokes, R. H. Electrolyte Solutions; Butterworths: London, 1959. (b) Stokes, R. H.; Robinson, R. A. J. Am. Chem. Soc. 1948, 70, 1870. (c) Bates, R. G.; Staples, B. R.; Robinson, R. A. Anal. Chem. 1970, 42, 867. (22) (a) Koryta, J.; Dvorak, J.; Bohackova, V. Electrochemistry; Methuen: London, 1970; p 32. (b) Brunh, H.; Nigan, S.; Holzwarth, J. F. Faraday Discuss. Chem. Soc. 1982, 74, 129. (23) Moya´, M. L.; Rodrı´guez, A.; Sa´nchez, F. Inorg. Chim. Acta 1992, 197, 227.

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